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Remembering General Chemistry: Electronic Structure and Bonding

Chapter 1. Remembering General Chemistry: Electronic Structure and Bonding. Paula Yurkanis Bruice University of California, Santa Barbara. What is Organic Chemistry?. Organic compounds: from living organisms (with a vital force) Inorganic compounds: from minerals

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Remembering General Chemistry: Electronic Structure and Bonding

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  1. Chapter 1 Remembering General Chemistry: Electronic Structure and Bonding Paula Yurkanis Bruice University of California, Santa Barbara

  2. What is Organic Chemistry? • Organic compounds: from living organisms (with a vital force) • Inorganic compounds: from minerals (with a vital force) organic chemistry = compounds that contain carbon

  3. What Makes Carbon So Special? • Atoms to the left of carbon give up electrons. • Atoms to the right of carbon accept electrons. • Carbon shares electrons.

  4. The Structure of an Atom Protons are positively charged. Neutrons have no charge. Electrons are negatively charged. Atomic number = # of protons Atomic number of carbon = 6 Neutral carbonhas six protons andsix electrons.

  5. The Distribution of Electrons in an Atom • The first shell is closest to the nucleus. • The closer the atomic orbital is to the nucleus, the lower its energy. • Within a shell, s < p.

  6. Relative Energies of the Atomic Orbitals

  7. -Aufbau principle: An electron goes into the atomic orbital with the lowest energy. -Pauli exclusion principle: No more than twoelectrons can be in an atomic orbital. -Hund’s rule: An electron goes into an emptydegenerateorbital rather than pairing up.

  8. Atoms on the Left Side of the Periodic Table Lose an Electron

  9. Atoms on the Right Side of the Periodic Table Gain an Electron

  10. A Hydrogen Atom Can Lose or Gain an Electron

  11. An Ionic Bond is Formed by the Attraction Between Ions of Opposite Charge

  12. Covalent Bonds are Formed by Sharing Electrons Nonpolar covalent bond = bonded atoms are the same Polar covalent bond = bonded atoms are different

  13. How Many Bonds Does an Atom Form? Lewis Structures

  14. Bond Polarity Depends on the Difference in Electronegativity

  15. Polar Covalent Bonds

  16. Dipole Moments the greater the difference in electronegativity, the greater the dipole moment, and the more polar the bond

  17. Electrostatic Potential Maps

  18. Formal Charge Formal Charge = the number or valence electrons– 1/2 the number of bonding electrons Formal Charge = the number or valence electrons– the number of bonds or

  19. Neutral Carbon Forms Four Bonds if carbon does not form four bonds, it has a charge (or it is a radical)

  20. Neutral Nitrogen Forms Three Bonds Nitrogen has one lone pair. If nitrogen does not form three bonds, it is charged.

  21. Neutral Oxygen Forms Two Bonds Oxygen has two lone pairs. If oxygen does not form two bonds, it is charged.

  22. Hydrogen and the Halogens Form One Bond A halogen has three lone pairs. if hydrogen or halogen does not form one bond, it has a charge (or it is a radical)

  23. The Number of Bonds and Lone Pairs • Halogen = 3 lone pairs • Oxygen = 2 lone pairs • Nitrogen = 1 lone pair

  24. Lewis Structures

  25. How to Draw a Lewis Structure NO3– Determine the total number of valence electrons (5 + 6 + 6 + 6 = 23). Because they are negatively charged, add another electron = 24. Avoid O–O bonds. Check for formal charges.

  26. Kekulé Structures and Condensed Structures

  27. What is an Atomic Orbital?

  28. A Standing Wave

  29. The Lobes of a p Orbital Have Opposite Phases

  30. The Three p Orbitals

  31. The Bonding in H2

  32. Waves Can Reinforce Each Other; Waves Can Cancel Each Other

  33. Atomic Orbitals Combine to Form Molecular Orbitals

  34. Side-to-Side Overlap of In-Phase p Orbitals Forms a π Bond

  35. The p Orbital of the More Electronegative Atom Contributes More to the Bonding MO

  36. The Bonding in Methane

  37. In Order to Form Four Bonds,Carbon Must Promote an Electron

  38. Four Orbitals are Mixed to Form Four Hybrid Orbitals An sp3 orbital has a large lobe and a small lobe.

  39. The Carbon in Methane is sp3 Carbon is tetrahedral. The tetrahedral bond angle is 109.5°.

  40. The Bonding in Ethane

  41. The Bonding in Ethane

  42. End-On Overlap of Orbitals Forms a (σ) Bond

  43. Bonding in Ethene

  44. An sp2 Carbon Has Three sp2 Orbitals and One p Orbital

  45. The Carbons in Ethene are sp2

  46. Bonding in Ethyne

  47. The Two sp Orbitals Point in Opposite Directions The Two p Orbitals are Perpendicular

  48. The Carbons in Ethyne are sp

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